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Chapter 5: Soap
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Introductory Activity
Fill a test tube with an inch of waterAdd a squirt of cooking oil to the test tube.
ObserveStopper, shake & observeAdd a few drops of soap. ObserveStopper, shake & observeWith another test tube, add water & soap only.
Observe.Compare the two test tubes.Make particle visualizations describing each test
tube.
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Introductory Activity
What ideas do you have about how soap works?
What kinds of things do advertising and marketing tell you?
What do the soap companies want you to know about how soap works?
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Soap
This chapter will introduce the chemistry needed to understand how soap worksSection 5.1: Types of bondsSection 5.2: Drawing MoleculesSection 5.3: Compounds in 3DSection 5.4: Polarity of MoleculesSection 5.5: Intermolecular ForcesSection 5.6: Intermolecular Forces and
Properties
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Soap
Inter-molecular forces
Inter-molecular forces
Works based on
Molecular Geometry
Molecular Geometry
Bonding types &
Structures
Bonding types &
Structures
Determined by
Determined by
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Section 5.1—Types of Bonds
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Why atoms bond
Atoms are most stable when they’re outer shell of electrons is full
Atoms bonds to fill this outer shellFor most atoms, this means having 8
electrons in their valence shellCalled the Octet Rule
Common exceptions are Hydrogen and Helium which can only hold 2 electrons.
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One way valence shells become full
Na-
-
- --
-
-- - -
Cl-
-
- --
-
-- - -
-
-
-
--
-
-
Sodium has 1 electron in it’s valence shell
Chlorine has 7 electrons in it’s valence shell
Some atoms give electrons away to reveal a full level underneath.
Some atoms gain electrons to fill their current valence shell.
-
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One way valence shells become full
Na-
-
- --
-
-- - -
Cl-
-
- --
-
-- - -
-
-
-
--
-
-
-+ -
The sodium now is a cation (positive charge) and the chlorine is now an anion (negative charge).
These opposite charges are now attracted, which is an ionic bond.
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Ionic Bonding—Metal + Non-metal
Metals have fewer valence electrons and much lower ionization energies (energy needed to remove an electron) than non-metals
Therefore, metals tend to lose their electrons and non-metals gain electrons
Metals become cations (positively charged)Non-metals become anions (negatively charged)The cation & anion are attracted because of
their charges—forming an ionic bond
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Bonding between non-metals
When two non-metals bond, neither one loses or gains electrons much more easily than the other one.
Therefore, they share electronsNon-metals that share electrons evenly
form non-polar covalent bondsNon-metals that share electrons un-evenly
form polar covalent bonds
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Metals bonding
Metals form a pool of electrons that they share together.
The electrons are free to move throughout the structure—like a sea of electrons
Atoms aren’t bonded to specific other atoms, but rather to the network as a whole
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Bond type affects properties
The type of bonding affects the properties of the substance.
There are always exceptions to these generalizations (especially for very small or very big molecules), but overall the pattern is correct
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Melting/Boiling Points
Ionic bonds tend to have very high melting/boiling points as it’s hard to pull apart those electrostatic attractionsThey’re found as solids under normal conditions
Polar covalent bonds have the next highest melting/boiling pointsMost are solids or liquids under normal conditions
Non-polar covalent bonds have lower melting/boiling pointsMost are found as liquids or gases
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Solubility in Water
Ionic & polar covalent compounds tend to be soluble in water
Non-polar & metallic compounds tend to be insoluble
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Conductivity of Electricity
In order to conduct electricity, charge must be able to move or flow
Metallic bonds have free-moving electrons—they can conduct electricity in solid and liquid state
Ionic bonds have free-floating ions when dissolved in water or in liquid form that allow them conduct electricity
Covalent bonds never have charges free to move and therefore cannot conduct electricity in any situation
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Section 5.2—Drawing Molecules
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Drawing Molecules on Paper
Lewis Structures (or Dot Structures) are one way we draw molecules on paper
Since paper is 2-D and molecules aren’t, it’s not a perfect way to represent how molecules bond…but it’s a good way to begin to visualize molecules
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Drawing Ionic Compounds
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1: How many valence electrons are in an atom?
The main groups of the periodic table each have 1 more valence electron than the group before it.
1 2 3 4 5 6 7 8
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2: Placing electrons around an atom
When atoms bond, they have 4 orbitals available (1 “s” and 3 “p”s). There are 4 places to put electrons
Put one in each spot before doubling up!
Example:Draw the
Lewis Structure for an oxygen
atom
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3: Transfer electrons in ionic bonding
Transfer electrons from metal atoms to non-metal atoms, keeping track of their new charge
Example:Draw the
Lewis Structure for
KCl
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4: Add more atoms if needed
If the transfer from one atom to another doesn’t result in full outer shells, add more atoms
Example:Draw the
Lewis Structure the
ionic compound of
Barium fluoride
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4: Add more atoms if needed
If the transfer from one atom to another doesn’t result in full outer shells, add more atoms
FBa
Barium has 2 electron
Fluorine has 7 electrons
Example:Draw the
Lewis Structure the
ionic compound of
Barium fluoride
F
Add another fluorine atom
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A note about Ionic Dot Structures
The atoms are not sharing the electrons—make sure you clearly draw the atoms separate!
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Drawing Covalent Compounds
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Tips for arranging atoms
Hydrogen & Halogens (F, Cl, Br, I) can only bond with one other atom—they can’t go in the middle of a molecules
Always put them around the outside
In general, write out the atoms in the same order as they appear in the chemical formula
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Repeat first two steps from before
1. Use the periodic table to decide how many electrons are around each atom
2. Write the electrons around each atom
Example:Draw the
Lewis Structure for
CH4
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H
H
Repeat first two steps from before
1. Use the periodic table to decide how many electrons are around each atom
2. Write the electrons around each atom
Example:Draw the
Lewis Structure for
CH4
Remember, “H” can’t go in the middle…put them around the Carbon!
C HH
Carbon has 4 electrons
Each hydrogen has 1
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H
H
3: Count electrons around each atom
Any electron that is being shared (between two atoms) gets to be counted by both atoms!
All atoms are full with 8 valence electrons (except H—can only hold 2)
Example:Draw the
Lewis Structure for
CH4
C HHCarbon has 8
Each Hydrogen has 2
All have full valence shells—drawing is correct!
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Bonding Pair
Pair of electrons shared by two atoms…they form the “bond”
H
HC HH
Bonding pair
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What if they’re not all full after that?
Sometimes, the first 3 steps don’t leave you with full valence shells for all atoms
Example:Draw the
Lewis Structure for
CH2O
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Double Bonds & Lone Pairs
Double bonds are when 2 pairs of electrons are shared between the same two atoms
Lone pairs are a pair of electrons not shared—only one atom “counts” them
HC OH
Double Bond
Lone pair
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And when a double bond isn’t enough…
Sometimes forming a double bond still isn’t enough to have all the valence shells full
Example:Draw the
Lewis Structure for
C2H2
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Properties of multiple bonds
Single Bond
Double Bond
Triple Bond
Shorter bonds (atoms closer together)
Stronger bonds (takes more energy to break)
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Polyatomic Ions
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Polyatomic Ions
They are a group of atoms bonded together that have an overall charge
Example:Draw the
Lewis Structure for
CO3-2
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Polyatomic Ions
They are a group of atoms bonded together that have an overall charge
Example:Draw the
Lewis Structure for
CO3-2
C O
Now the Carbon and the one oxygen have 8…but the other two oxygen atoms still only have 7
OO
This is a polyatomic ion with a charge of “-2”…that means we get to “add” 2 electrons!
-2
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Covalent bond within…ionic bond between
Polyatomic ions have a covalent bond within themselves…
But an ionic bond with other ions
Covalent bonds within
Ionic bond with other ions
C OOO
-2
Na
Na
+1
+1
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Isomers
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More than one possibility
Often, there’s more than one way to correctly draw a Dot Structure
HC CH CHH
HC CH CH
H
Chemical Formula: C3H4
Chemical Formula: C3H4
Contains 2 sets of double bonds between carbons
Contains 1 triple bond and 1 single bond between carbons
Both structures have full valence shells!
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Both are “correct”
The chemical formula alone does not give you enough information to differentiate between the two structures
HC CH CHH
HC CH CH
H
Chemical Formula: C3H4
You’ll learn in Chapter 11 how to differentiate between these two structures
with chemical names
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Isomers
Isomers: Structures with the same chemical formula but different chemical structure
Atoms must be bonded differently (multiple versus single bonds) or in a different order) but have the same overall chemical formula to be isomeric structures
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Section 5.3—Molecules in 3D
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Bonds repel each other
Bonds are electrons. Electrons are negatively charged
Negative charges repel other negative charges
Bonds repel each other
Molecules arrange themselves in 3-D so that the bonds are as far apart as possible
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ValenceShellElectronPairRepulsionTheory
Valence Shell Electron Pair Repulsion Theory (VSEPR Theory)
Outer shell of electrons involved in bonding
Bonds are made of electron pairs
Those electron pairs repel each other
Attempts to explain behavior
This theory (that bonds repel each other because they’re like charges) attempts to explain why molecules form the shapes they form
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What shapes do molecules form?
Linear
2 bonds, no lone pairs
Trigonal planar3 bonds, no lone pairs
Indicates a bond coming out at you
Indicates a bond going away from you
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What shapes do molecules form?
Tetrahedron
4 bonds, no lone pairs
Trigonal bipyramidal
5 bonds, no lone pairs
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What shapes do molecules form?
Octahedron
6 bonds, no lone pairs
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Lone Pairs
Lone pairs are electrons, too…they must be taken into account when determining molecule shape since they repel the other bonds as well.
But only take into account lone pairs around the CENTRAL atom, not the outside atoms!
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What shapes do molecules form?
Bent
2 bonds, 1 lone pair
Trigonal pyramidal
3 bonds, 1 lone pair
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What shapes do molecules form?
Bent
2 bonds, 2 lone pairs
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Lone Pairs take up more space
Lone pairs aren’t “controlled” by a nucleus (positive charge) on both sides, but only on one side.
This means they “spread out” more than a bonding pair.
They distort the angle of the molecule’s bonds away from the lone pair.
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109.5°
C
105°
O
Example of angle distortion
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Ionic Compound structures
Ionic compounds are made of positive and negative ions.
They pack together so that the like-charge repulsions are minimized while the opposite-charge attractions are enhanced.
Na+1 Cl-1
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Section 5.4—Polarity of Molecules
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Electronegativity
The pull an atom has for the electrons it shares with another atom in a bond.
Electronegativity is a periodic trendAs atomic radius increases and number of
electron shells increases, the nucleus of an atom has less of a pull on its outermost electrons
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Periodic Table with Electronegativies
increases
decreases
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Polar Bond
A polar covalent bond is when there is a partial separation of charge
One atom pulls the electrons closer to itself and has a partial negative charge.
The atom that has the electrons farther away has a partial positive charge
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Two atoms sharing equally
N N
Each nitrogen atom has an electronegativity of 3.0
They pull evenly on the shared electrons
The electrons are not closer to one or the other of the atoms
This is a non-polar covalent bond
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Atoms sharing almost equally
Electronegativities: H = 2.1 C = 2.5
The carbon pulls on the electrons slightly more, pulling them slightly towards the carbon
Put the difference isn’t enough to create a polar bond
This is a non-polar covalent bond
C HH
H
H
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Sharing unevenly
Electronegativities: H = 2.1 C = 2.5 O = 3.5
The carbon-hydrogen difference isn’t great enough to create partial charges
But the oxygen atoms pulls significantly harder on the electrons than the carbon does. This does create a polar covalent bond
This is a polar covalent bond
C OH
H
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Showing Partial Charges
There are two ways to show the partial separation of chargesUse of “” for “partial” Use of an arrow pointing towards the partial
negative atom with a “plus” tail at the partial positive atom
C OH
H
+ -C OH
H
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Ionic Bonds
Ionic bonds occur when the electronegativies of two atoms are so different that they can’t even share unevenly…one atom just takes them from the other
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How to determine bond type
Find the electronegativies of the two atoms in the bond
Find the absolute value of the difference of their valuesIf the difference is 0.4 or less, it’s a non-polar
covalent bondIf the difference is greater than 0.4 but less than
1.4, it’s a polar covalent bondIf the difference is greater than 1.4, it’s an ionic
bond
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Let’s Practice
Example:If the bond
is polar, draw the polarity arrow
C – H
O—Cl
F—F
C—Cl
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Polar Bonds versus Polar Molecules
Not every molecule with a polar bond is polar itselfIf the polar bonds cancel out then the molecule
is overall non-polar.
The polar bonds cancel out.No net dipole
The polar bonds do not cancel out.
Net dipole
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The Importance of VSEPR
You must think about a molecule in 3-D (according to VSEPR theory) to determine if it is polar or not!
Water drawn this way shows all the polar bonds canceling out. But water drawn in
the correct VSEPR structure, bent, shows the polar bonds don’t cancel out!
Net dipole
H O H
O H H
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Let’s Practice
Example:Is NH3 a
polar molecule?
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Section 5.5—Intermolecular Forces
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Intra- versus Inter-molecular Forces
So far this chapter has been discussing intramolecular forcesIntramolecular forces = forces within the
molecule (chemical bonds)
Now let’s talk about intermolecular forcesIntermolecular forces = forces between
separate molecules
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Breaking Intramolecular forces
Breaking of intramolecular forces (within the molecule) is a chemical change2 H2 + O2 2 H2O
Bonds are broken within the molecules and new bonds are formed to form new molecules
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Breaking Intermolecular forces
Breaking of intermolecular forces (between separate molecules) is a physical changeBreaking glass is breaking the intermolecular
connections between the glass molecules to separate it into multiple pieces.
Boiling water is breaking the intermolecular forces in liquid water to allow the molecules to separate and be individual gas molecules.
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London Dispersion Forces
All molecules have electrons.
Electrons move around the nuclei. They could momentarily all “gang up” on one side
This lop-sidedness of electrons creates a partial negative charge in one area and a partial positive charge in another.
+ Positively charged nucleus - Negatively charged electron
+-
-
-
-
Electrons are fairly evenly dispersed.
+--
- -As electrons move, they “gang up” on one side.
+
-
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London Dispersion Forces
Once the electrons have “ganged up” and created a partial separation of charges, the molecule is now temporarily polar.
The positive area of one temporarily polar molecule can be attracted to the negative area of another molecule.
+ - + -
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Strength of London Dispersion Forces
Electrons can gang-up and cause a non-polar molecule to be temporarily polar
The electrons will move again, returning the molecule back to non-polar
The polarity was temporary, therefore the molecule cannot always form LDF.
London Dispersion Forces are the weakest of the intermolecular forces because molecules can’t form it all the time.
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Strength of London Dispersion Forces
Larger molecules have more electrons
The more electrons that gang-up, the larger the partial negative charge.
The larger the molecule, the stronger the London Dispersion Forces
Larger molecules have stronger London Dispersion Forces than smaller molecules.
All molecules have electrons…all molecules can have London Dispersion Forces
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Dipole Forces
Polar molecules have permanent partial separation of charge.
The positive area of one polar molecule can be attracted to the negative area of another molecule.
+ - + -
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Strength of Dipole Forces
Polar molecules always have a partial separation of charge.
Polar molecules always have the ability to form attractions with opposite charges
Dipole forces are stronger than London Dispersion Forces
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Hydrogen Bonding
Hydrogen has 1 proton and 1 electron.There are no “inner” electrons. It bonds with the only
one it has.When that electron is shared unevenly (a polar
bond) with another atom, the electron is farther from the hydrogen proton than usual.This happens when Hydrogen bonds with Nitrogen,
Oxygen or FluorineThis creates a very strong dipole (separation of
charges) since there’s no other electrons around the hydrogen proton to counter-act the proton’s positive charge.
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Strength of Hydrogen Bond
Hydrogen has no inner electrons to counter-act the proton’s charge
It’s an extreme example of polar bonding with the hydrogen having a large positive charge.
This very positively-charged hydrogen is highly attracted to a lone pair of electrons on another atom.
This is the strongest of all the intermolecular forces.
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Hydrogen Bond
N
H H
N
H H
Hydrogen bond
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Section 5.6—Intermolecular Forces & Properties
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IMF’s and Properties
IMF’s are Intermolecular ForcesLondon Dispersion ForcesDipole interactionsHydrogen bonding
The number and strength of the intermolecular forces affect the properties of the substance.
It takes energy to break IMF’sEnergy is released when new IMF’s are
formed
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IMF’s and Changes in State
Some IMF’s are broken to go from solid liquid. All the rest are broken to go from liquid gas.
Breaking IMF’s requires energy.
The stronger the IMF’s, the more energy is required to melt, evaporate or boil.
The stronger the IMF’s are, the higher the melting and boiling point
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Water
Water is a very small moleculeIn general small molecules have low melting and
boiling pointsBased on it’s size, water should be a gas under
normal conditionsHowever, because water is polar and can form
dipole interactions and hydrogen bonding, it’s melting point is much higher
This is very important because we need liquid water to exist!
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IMF’s and Viscosity
Viscosity is the resistance to flowMolasses is much more viscous than
water
Larger molecules and molecules with high IMF’s become inter-twined and “stick” together more
The more the molecules “stick” together, the higher the viscosity
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Solubility
In order from something to be dissolved, the solute and solvent must break the IMF’s they form within itself
They must then form new IMF’s with each other
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Solubility
- +
- +
- + - +- +
Solvent, water (polar)
+
-
- + Solute, sugar (polar)
Water particles break some intermolecular forces with other water molecules (to allow them to spread out) and begin to form new ones with the sugar molecules.
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Solubility
Solvent, water (polar)
+
-
- + Solute, sugar (polar)
As new IMF’s are formed, the solvent “carries off” the solute—this is “dissolving”
- +
- +
- +- + - +
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Solubility
If the energy needed to break old IMF’s is much greater than the energy released when the new ones are formed, the process won’t occurAn exception to this is if more energy is added
somehow (such as heating)
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Oil & Water
Water has London Dispersion, Dipole and hydrogen bonding. That takes a lot of energy to break
Water can only form London Dispersion with the oil. That doesn’t release much energy
Much more energy is required to break apart the water than is released when water and oil combine.
Water is polar and can hydrogen bond, Oil is non-polar.
Therefore, oil and water don’t mix!
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Surface Tension
Surface tension is the resistance of a liquid to spread out.This is seen with water on a freshly waxed car
The higher the IMF’s in the liquid, the more the molecules “stick” together.
The more the molecules “stick” together, the less they want to spread out.
The higher the IMF’s, the higher the surface tension.
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Soap & Water
Soap has a polar head with a non-polar tail
The polar portion can interact with water (polar) and the non-polar portion can interact with the dirt and grease (non-polar).
Polar head
Non-polar tailSoap
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Soap & Water
The soap surrounds the “dirt” and the outside of the this Micelle can interact with the water.
The water now doesn’t “see” the non-polar dirt.
Dirt
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Soap & Surface Tension
The soap disturbs the water molecules’ ability to form IMF’s and “stick” together.
This means that the surface tension of water is lower when soap is added.
The lower surface tension allows the water to spread over the dirty dishes.
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What did you learn about soap?
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Soap
Inter-molecular forces
Inter-molecular forces
Works based on
Molecular Geometry
Molecular Geometry
Bonding types &
Structures
Bonding types &
Structures
Determined by
Determined by